GASES

1. GASES

2. KINETIC MOLECULAR • THEORY OF GASES: • This theory is used to describe the nature • of gases and establish the gas laws.

3. KINETIC MOLECULAR • THEORY OF GASES: • Gases consist of small particles in rapid, • random motion. • 2) There is very little attraction between • gas particles. • 3) Average kinetic energy is proportional • to Kelvin temperature. • 4) Gas particles are in constant, straight-line • motion.

5. PRESSURE: • As gas particles hit the walls of their • container, they push on it. This force is • the pressure. • Pressure (P) = force/area

6. ATMOSPHERIC PRESSURE: • This is the pressure on the air around us. • Atmospheric pressure decreases as • elevation increases. • Atmospheric pressure is measured with a • BAROMETER, so it is often called barometric • pressure. • Don’t confuse ATMOSPHERIC PRESSURE • with the pressure unit ATMOSPHERE, which • is 760 mmHg.

8. Evangelista Torricelli

9. BOYLE’S LAW: • Pressure and volume are inversely • proportional to each other. This means, • when pressure increases, volume decreases • and, when pressure decreases, volume • increases. • At constant n and T, PV = constant • or • P1V1 = P2V2.

10. Robert Boyle (It was common for men to wear wigs back then)

11. BOYLE’S LAW: • As volume gets smaller, the particles hit • the walls more often, raising the pressure.

12. BOYLE’S LAW: • A sample of helium has 4.0 L and 705 mm Hg. What is the volume of the sample if the pressure is changed to 745 mm Hg?

13. CHARLES’ LAW: • Volume is directly proportional to • temperature. That is, when temperature • increases, volume increases and when • temperature decreases, volume decreases. • At constant n and P, V/T = constant • or • V1/T1 = V2/T2.

14. CHARLES’ LAW: • When temperature increases, gas particles • move faster. This means they hit the walls • of the container more often, causing it • to expand.

15. Jacque Charles Invented the hydrogen-filled balloon and on December 1, 1783, he ascended into the air and became possibly the first man in history to witness a double sunset.

16. CHARLES’ LAW • A balloon filled with air has a volume of • 2.5 L. What is the volume of the balloon • if you bring it inside (25 °C) from the cold • weather (-5 °C)? • Remember, you MUST use Kelvin: • 25 °C = 25 °C + 273 = 298 K • -5 °C = -5 °C + 273 = 268 K

17. GAY-LUSSAC’S LAW: • Pressure is directly proportional to Kelvin • temperature. That is, as temperature • increases, pressure increases and, as • temperature decreases, pressure decreases. • At constant n and V, P/T = constant • or • P1/T1 = P2/T2.

18. GAY-LUSSAC’S LAW: • As temperature increases, gas particles • move faster. This means that they hit the • walls of the container more often, increasing • the force applied to the walls.

19. Joseph Louis Guy-Lussac Actually discovered Charles’ Law, too. Also a balloonist…on Sept. 16, 1804, rose to just over 23,000 feet (about 4.3 miles), a world altitude record for over 50 years.

20. Why should you • study chemistry? • The same year a 23-year-old Gay-Lussac discovered this law, he had occasion to walk into a linen draper's shop in Paris and there he made a wonderous discovery. He found the 17-year-old shopgirl reading a chemistry textbook while • waiting for customers. Needless to say, he was intrigued by • this and made more visits to the shop. In 1808, he and • Josephine were married and over the years, five little • Gay-Lussacs were added to the scene.

21. GAY-LUSSAC’S LAW: • You fill your car tires to 35 psi in the • summer when it’s 31 °C. At Thanksgiving, • when it’s -7 °C, you notice that your tire is • low. Why? • Remember, you MUST use Kelvin: • 31 °C = 31 °C + 273 = 304 K • -7 °C = -7 °C + 273 = 266 K

22. VAPOR PRESSURE • AND BOILING POINT: • The pressure of the vapor above a liquid • is called the VAPOR PRESSURE. • When the vapor pressure equals the • external (atmospheric) pressure, that is • the BOILING POINT.

25. VAPOR PRESSURE • AND BOILING POINT: • As elevation increases, pressure decreases. • This is why water boils at lower temperatures • at high altitudes.

26. COMBINED GAS LAW: • We can take all of the pressure-volume- • temperature relationships we’ve just • talked about and combine them into • one gas law: • (P1V1)/T1 = (P2V2)/T2 • If you remember this gas law, you can • figure out the others.

27. COMBINED GAS LAW: • A scuba diver 40 ft below the ocean surface • inhales 50.0 mL of compressed air in a scuba • tank at a pressure of 3.00 atm and a • temperature of 8 °C. What is the pressure • of air in the lungs if the gas expands to • 150.0 mL at a body temperature of 37 °C? • Remember, you MUST use Kelvin: • 8 °C = 8 °C + 273 = 281 K • 37 °C = 37 °C + 273 = 310 K

29. AVOGADRO’S LAW: • In each of the previous gas laws, we’ve • assumed that the amount (grams or moles) • of gas doesn’t change. • Gas amount and volume are directly • proportional to each other. This means, • when you have more gas, volume increases • and, when you have less gas, volume • decreases. • V1/n1 = V2/n2

30. AVOGADRO’S LAW: • The more gas particles there are, the more • will hit the walls of the container, causing • it to expand.

31. STP AND MOLAR VOLUME: • Avogadro’s law tells us that two gases at • the same temperature and pressure will • have the same volume. • STANDARD TEMPERATURE AND PRESSURE (STP) is often used when comparing gases. • STANDARD TEMPERATURE = 0 °C (273 K) • STANDARD PRESSURE is 1 atm (760 mm Hg)

32. STP AND MOLAR VOLUME: • At STP, the volume of one mole of ANY gas • is 22.4 L. • Conversion • Factors: • 1 mol @STP • 22.4 L • or • ___22.4 L___ • 1 mol @ STP

33. IDEAL GAS LAW: • When dealing with a single gas, but no • changes in conditions, we can describe • the gas with the IDEAL GAS LAW: • PV = nRT. • R is a constant, 0.0821 L x atm/mole x K.

34. IDEAL GAS LAW: • You fill your bike tire with 3.0 L of air at • 7.15 atm on a summer day that is 30 °C. • How many moles of air are in your bike • tire? • Remember, you MUST use Kelvin: • 30 °C = 30 °C + 273 = 303 K

35. REMEMBER THIS: Moles A  Moles B Grams of A Grams of B x g/mole of B x mole/g of A

36. Moles A  Moles B Grams of A Grams of B x g/mole of B x mole/g of A n = PV/RT PV = nRT Gas P, V, or T of B

37. Problem 8.98 on Page 275: Hydrogen gas can be produced in the laboratory though the reaction of magnesium metal with hydrochloric acid: Mg (s) + 2 HCl  MgCl2 (aq) + H2 (g) What is the volume, in liters, of H2 gas produced at 24 °C and 835 mm Hg, from the reaction of 12.0 g of Mg?

38. DALTON’S LAW: • The pressure of a mixture of gases is equal • to the sum of the pressure of each of the • gases. • The total pressure of a gas mixture depends • on the total number of gas particles, not on • the types of particles. • The pressure of each gas is equal to the • total pressure times the (molar) percent • of the gas in the mixture. • Ptot = P1 + P2 + P3 + …

39. TOTALPRESSURE: • For example, at STP, one mole of gas particles in a volume of 22.4 L will exert the same pressure as one mole of a mixture of gas particles in 22.4 L. 1.0 mole N2 0.4 mole O2 0.6 mole He 1.0 mole 0.5 mole O2 0.3 mole He 0.2 mole Ar 1.0 mole 1.0 atm 1.0 atm 1.0 atm

42. DALTON’S LAW: • During space walks, astronauts breathe • a mixture of nitrogen and oxygen that has • a total pressure of 7.0 atm. The partial • pressure of nitrogen in this mixture is • 5.6 atm. What is the partial pressure of • oxygen? What is the composition of this • mixture?

43. DALTON’S LAW: • Dalton’s Law is especially important for • breathing. Not only for things like scuba • tanks, but especially in medical situations. • This is why “blood gases” are often • measured.

44. When the lungs expand, the pressure in the lungs decreases. • Inhalation occurs as air flows towards the lower pressure in the lungs.

45. When the lung volume decreases, pressure within the lungs increases. • Exhalation occurs as air flows from the higher pressure in the lungs to the outside.

46. (c) Douglas E. Raynie, South Dakota State University, 2004

47. (c) Douglas E. Raynie, South Dakota State University, 2004